Bonding and Structure

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Atomic Structure & Chemical Bonding Notes

(Credits to Mustafa Asif)

Chemical Bonding Worksheet --> Answers

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Summary

Nuclear model of an atom

• • Proton number: number of protons in an atom

  • Nucleon number: number of protons + number of neutrons

• Number of electrons = number of protons

  • Isotopes: atoms of the same element with same proton number but different number of neutrons

    • isotopes have slightly different physical properties but have identical chemical properties as they have the same number and arrangement of electrons

  • Valence electrons: the electrons in the outermost shell

    • used to form chemical bonds

Electronic Configuration

• • Electrons are placed in orbits.

  • First shell contains maximum 2 electrons. Second shell and so and so for has maximum of 8 electrons.

  • To write electronic configuration we write as n.n.n.... where first n denotes the first shell, second the second shell and so and so for.

    • E.g. Sulfur has electronic configuration of 2.8.6

  • The valence electrons is the number of electrons of the outermost shell.

    • Sulphur has 6 valence electrons.

Relation with Periodic Table

• • Elements in same horizontal row: Period

  • Elements in same vertical column: Group

    • Group 1 has 1 valency

    • Group 2 has 2 valency

    • Group 3 has 3 valency and so on.

    • Group 0 has full valency which makes it having stable electronic configuration.

  • Down the group the number of shells increases.

Structure of Matter

a. Atom: Smallest particle of an element

b. Molecule: group of two or more atoms chemically joined together, e.g. chlorine molecule has 2 chlorine atoms

c. Element: Element is a substance that cannot be broken down into simpler substances by chemical nor physical methods.

Classifying elements

• • Classifying by state. E.g. some elements are solids, some liquids, some gases.

  • Classifying by metals and non-metals. E.g. most elements are metals, semi-metals are metalloids (having properties of metals & non-metals), some are non-metals

  • Classifying by periodicity.From left-right elements change from metal to non-metal

d. Compound: substance containing 2 or more elements chemically joined together

e.g. Magnesium is an element; oxygen is an element – they can only be burnt to form magnesium oxide compound.

Composition of compounds

• • Ions or molecules make up compounds

  • Ions are atoms having electrical charge

  • E.g. NaCl made up of 2 ions; positively charged Na, negatively charged Cl.

e. Mixtures: contains 2 or more substances not chemically joined together. e.g. seawater is made up of water and NaCl (salt); oxygen in air varies

Compounds vs Mixtures

Ionic bond

• • Ionic bonding is the transfer of electrons from one atom to another to become achieve an inert gas configuration, forming ions.

  • Ionic bonds are formed between METALLIC and NON- METALLIC ATOMS ONLY.

  • Metals lose electrons to form positive ions (cations)

  • Non-metals gain electrons to form negative ions (anions)

  • The formation of ions is resulted from transfer of atoms from one atom to another atom(s), which the ions produced are of opposite charges, and unlike charges attract, causing them to be held together with a strong force.

Example: Formation of NaCl

• • Sodium atom loses an electron by transferring the electron to chlorine atom, making both stable.

  • The loss of electron forms cation, Na+, and the gain of electron forms anion, Cl-.

  • The opposite charges acquired by both ions attract to each other, forming a strong ionic bond of NaCl.

Deducing formula ionic compounds

• • We can know the charge of elements by looking at groups of periodic table.

    • Group I to group III elements has charge of +1, increasing to +3, going to the right.

    • Group V to group VII elements has charge of -3, decreasing to -1, going to the right.

E.g. Aluminium sulfate

We have to balance the charges to make a stable bond.

Ions present:

Al³⁺ SO₄^2-

Al³⁺ SO₄^2-

SO₄^2-

Total change: 6+ 6-

Therefore, the formula is Al₂(SO₄)₃

1. The symbol of metal ion should always be first, e.g. NaCl

2. Polyatomic ion should be placed in brackets, e.g. Fe(NO₃)₂

Structure and Properties of ionic bond

Structure

• • Ionic substances appear as giant lattice structures which the ions are held together by electrostatic force between oppositely charged ions.

  • To find the formula of ionic bond, say sodium chloride bond, by looking at lattice structure, we count the ratio of amount of metal ions to non-metal ions.

    • E.g. in sodium chloride, the ratio Na:Cl is 1:1, therefore the ionic formula is NaCl.

Properties

• • Ionic compounds are hard crystalline solids with flat sides and regular shapes because the ions are arranged in straight rows in strong ionic bonds.

  • Ionic compounds have very high melting points and boiling points.

  • The strong forces holding ionic compounds prevents them to evaporate easily. Hence, ionic compounds have no smell.

  • Solid ionic compounds don’t conduct electricity but they do when they are aqueous or molten.

    • This is because in liquid/aqueous state the ions which conduct electricity are free to move.

    • In solids, these ions are fixed in place.

  • Ionic compounds are soluble in water but insoluble in organic compounds.

    • This is because the ions attract water molecules which distrupts the crystal structure, causing them separate & go into solution.

    • Vice versa when in organic solvent.

Covalent bond

• • Covalent bonding is the sharing a pair of electrons to gain electronic configuration of an inert gas, usually for molecules.

  • Covalent bonds occur between NON-METALLIC ATOMS ONLY.

  • In covalent bond, we try to substitute the short of electrons of two or more atoms between each other to form the 2 or 8 valence electrons.

  • The shared electrons appear in pairs

Example: H₂ molecule

• • Hydrogen atom has one valency.

  • To become stable with hydrogen atom, it needs one more electron, just like helium which has 2 valency.

  • When 2 hydrogen atoms join, they share their electrons, on which, the share becomes 2 electrons, which is now a noble gas configuration, being shared between these 2 atoms.

  • We write the bond as H – H single bond, which means they share an electron pair (2 electrons).

Example: Cl₂ molecule

• • Cl atom has 7 valency and needs one electron, each, to form a noble gas configuration between two Cl atoms.

  • Hence they share an electron EACH to hence share 2 electrons between the atoms.

  • Hence, each Cl atom now has 8 valency which is a noble gas configuration.

Example: O₂ molecule

• • An O atom has 6 valency and needs 2 electrons, each, to form a noble gas configuration.

  • Hence, EACH SHARE THE AMOUNT OF ELECTRONS EACH SHORT OF, in this case – 2 electrons, to form stable molecule.

  • The contribution hence now become 4 electrons and what left on each oxygen atom is 4 electrons.

  • We combine each 4 electrons on oxygen atom with the 4 electrons shared and hence we get 8 valency for each oxygen atom – a noble gas configuration!

Example: H₂O molecule

• • Apart from oxygen sharing between oxygen atoms, it can have electrons with other atoms.

  • Oxygen needs 2 electrons and when bonded with hydrogen, which need an atom each, they combine to provide 2 electrons on both sides of oxygen bonded with hydrogen atoms.

  • Each hydrogen with oxygen atom form a single bond: O – H.

Example: CO₂ molecule

• • Carbon needs 4, oxygen needs 2.

  • We share two from oxygen part, WHICH HAS THE SMALLEST NUMBER OF SHORT ELECTRONS, TO SHARE THE AMOUNT OF ELECTRONS THAT ATOM NEEDS, to form 4 shared atoms.

  • Now oxygen is stable but carbon needs 2 more, which we now know they can get from another oxygen atom.

  • The atoms are now stable and since each bond has 2 pairs of electrons, we call this double bond: C = O.

Structure and Properties of covalent bond

Structure

a. Giant Covalent Bond

• Diamond

  • • Diamond has carbon atoms bonded with another carbon atoms in a tetrahedral arrangement which each carbon atom uses all its valence electrons to form 4 single covalent bonds with other 4 carbon atoms.

• Silicon Dioxide

  • • Silicon dioxide, SiO2, has silicon atoms bonded with another oxygen atoms in a tetrahedral arrangement which each silicon atom uses all its valence electrons to form 4 single covalent bonds with other 4 oxygen atoms.

• Graphite

  • • Graphite has flat layers of carbon atoms bonded strongly in hexagonal arrangement in which the layers are bonded to each other weakly.

Properties

• • It is a hard solid because it consists of many strong covalent bonds between atoms. This property makes it suitable as abrasives.

  • It has very high melting and boiling points.

  • It does not conduct electricity (except graphite) because there are no free electrons in covalent bonds since they are used to form bonds; hence electrons are in fixed positions. To conduct electricity, there must be free electrons.

  • All covalent structures are insoluble in water.

Metallic bond

• • Metallic bonding is bonding within atoms of metals caused by attractive force between positively charged metal ions and negatively charged free electrons.

  • The atoms are packed closely together in giant lattice structures.

BOND FORMING

• • Each atom in metal gives up valence electrons to form positive ions.

  • There are free electrons moving between the spaces and positive metal ions are attracted to the sea of electrons which hold the atoms together.

Structure and Properties of metallic bond

• • Metals can be bent (ductile) and can be stretched (malleable) because the layers of atoms in metals slide over each other when force is applied but will not break due to attractive force between electrons and metal ions.

  • Metals conduct electricity as it has free electrons which carries current.

  • Metals conduct heat as it has free electrons which gains energy when heated and moves faster to collide with metal atoms, releasing heat in collisions.

  • Metals have high melting and boiling points because the bonds between metals is very strong. Hence very high heat energy needed to break the bonds.

Simple Molecular Substances

• • Simple molecular substances are usually liquids/gases at r.t.p. because the molecules are not tightly bonded like in solids, hence free to move.

  • They have low melting and boiling points because the force of attraction is weak that they can be easily broken by heat.

  • Since they have low boiling points, they evaporate easily.

  • They don’t conduct electricity because they don’t have free electrons/ions which helps to conduct electricity.

  • Most of these are insoluble in water but soluble in organic solvent.

Noble gases

- Noble Gases like helium and argon are unreactive.

- do not form chemical bonds with other atoms as they have a stable electronic configuration.

- Atoms of other elements form chemical bonds so that they attain the electronic configuration of a noble gas

Drawing Structures

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